Chapter 3

Question 1

The Bohr Rutherford Model of the atoms is the

Answer 1

Planetary model

Question 2

Rutherford Contribution

Answer 2

• electrons move in space RANDOMLY around the nucleus
• the electrons account for the volume of the atom, and the nucleus accounts for the mass of an atom (neutrons + protons= atomic mass)
  –> Rutherford didn’t know about the neutrons in the nucleus
  –> It was his student, Chadwick, who later discovered the neutrons in the nucleus

Question 3

Bohr Contribution

Answer 3

• Electrons can exist only in a series of “allowed” energy levels or shells of “fixed” energies
• Electrons have to be in a shell, it cannot be “in-between” shells
• Each allowed energy state is given a quantum #: n=1,2,3…
• Lower shells have lower energies. Higher shells have higher energies

= planetary model

Question 4

If the electron is found in the lowest possible energy level, it is in its ___________________

Answer 4

ground state

Question 5

To jump from a lower level to a higher level, an electron must ____________________. The electron is said to be in an ________________.

Answer 5

absorb energy, excited state

Question 6

To fall back to a lower level from a higher level, an electron will …

Answer 6

release/emit the same amount of energy it absorbed. This energy may be observed as light

Question 7

Which absorption would require more energy?
From level 1-2 OR level 4-5?

Answer 7

From level 1-2
- Rmr the distance between shells decreases as # of shells get greater.

Question 8

Which absorption would release less energy?
From level 3-1 OR level 6-4?

Answer 8

level 6-4

Question 9

Electromagnetic Spectrum

Answer 9

• consists of light waves of diff frequencies & wavelengths

Question 10

Wavelength

Answer 10

λ - lamda
- the distance between any 2 corresponding points on adjacent waves –> crest-to-crest OR trough-to-trough

Question 11

Frequency

Answer 11

ν- nu
- the # of waves that pass in a specific amount of time at a certain point (per second)
- per second= 1/s= s^-1=hertz

Question 12

Wavelength & frequency are _______________________ to each other

Answer 12

inversely related
- As wavelength increases, frequency decreases
- As wavelength decreases, frequency increases

Question 13

electromagnetive spectrum from longest to shortest

Answer 13

Radio waves, microwaves, infrared, visible light, ultraviolet, x-rays, gamma

• visible light is obviously safe (i think)
• the first 3 are safe for use to be around
• the last 3 are not safe for us
• x-rays & gamma r not naturally found on earth

Question 14

Do shorter wavelengths have higher or lower energy?

Answer 14

Shorter wavelengths= higher energy
longer wavelengths= lower energy

Question 15

The visible spectrum is

Answer 15

the band of light waves (380-750) that the human can detect.
- 1m = 1,000,000,000 (1x10^9)nm
- colours of visible spectrum from smallest to largest wavelength= V,I,B,G,Y,O,R
- colours of visible spectrum from least to most energetic energy= R,O,Y,G,B,I,V

Question 16

A quantum is

Answer 16

fixed amount/value of energy
plural=quanta
- energy is quantized or restricted to specific levels –> think of ladder

Question 17

Emission Spectrum

Answer 17

WRITTEN ONE: an excited e- falls from a higher level to a lower one, releasing fixed amount of energy

TYPED ONE: it is the spectrum of frequency of EM (google: electromagnetic) radiation due to an e- making a transition from higher energy state to a lower energy state

• each element in the periodic table has its own “ladder of energies” or “emission line spectrum” signature

Question 18

How can we explain the lines observed for excited hydrogen

Answer 18

• when an e- gains a fixed amount of energy, it jumps from a higher level & becomes excited
• when this e- falls back to a lower level, a fixed amount of energy is released as light

Question 19

Hydrogen emission series

Answer 19

LYMAN: n* –> n=1, UV, not visible to naked eye
BALMER: n* –> n=2, visible, visible to naked eye
PASCHEN: n* –> n=3, IR, not visible to naked eye
BRAKET: n* –> n=4, IR, not visible to naked eye

• RMR ladder analogy with its gaps.
• smaller gap in each series= greater relative wavelengh
  –> shorter wavelength=greater energy, greater frequency

Question 20

Orbits VS Orbitals

Answer 20

ORBITS - e- travel around the nucleus in circular 2D pathway. The distances from the nucleus are fixed. No evidence exists for this model & Bohr’s atomic model of e- is abandoned

ORBITALS- the shape of an orbital is defined by the motion of the e- in that orbital. An orbital does not tell us the exact location of an e- where it is most likely to be found, only the 3D space where a 95% probability of finding an e- in it. Each orbital can hold a max of 2e-. (UNCERTAINTY PRINCIPLE) For diff allowed energy states, diff # & types of orbitals exist which are described by a set of quantum #

Question 21

The principal quantum # - n

Answer 21

• the energy levels
• shows how far an e- is from the nucleus (relative size of the nucleus
• the higher the #, the greater the energy of the e-

Question 22

Angular Momentum Quantum # - l

Answer 22

uses a # to represent the types & shapes of orbitals within each quantum level
l= 0 to n-1

Question 23

l=0

Answer 23

• an “s” spherical orbital
• the cloud’s density is not uniform throughout but is greater near the nucleus & decreases as we move away
• only one “s” orbital in each shell & can hold 2e-

Question 24

l=1

Answer 24

• 3 “p” perpendicular orbitals –> px, py, pz
• dumbbell shape= the lobes of each p orbital disappear at the origin where the nucleus is located –> meaning possibility of finding an e- at the nucleus is zero
• 3 “p” orbitals are = equal in energy (degenerate) & can hold max of 6 e-

Question 25

l=2

Answer 25

• 5 “d” diffuse orbitals
• 5 diff degenerate orbitals in diff direction
• can hold a max of 10e-

Question 26

l=3

Answer 26

• 7 “f” fundamental orbitals
• 7 diff degenerate orbitals in diff direction
• can hold a max of 14e-

Question 27

The Magnetic Quantum # -ml

Answer 27

• indicates the 3D orientation of an e- (which plane the e- is found in)
  -l <= ml<= l

Question 28

The Magnetic Spin Quantum # - ms

Answer 28

• the direction of e- spin
• each orbital can hold a max of 2 e- with opposite spin ( the Pauli Exclusion Principle)

Question 29

e- configurations using the condensed format

Answer 29

An atom’s full e- shells are represent by the symbol of the nearest preceding Noble Gas, in square brackets, followed by the e- configurations for the remaining e-

Question 30

Aufbau Principle

Answer 30

e- will fill lowest energy levels before higher ones are filled

Question 31

Pauli Exclusion Principle

Answer 31

each orbital holds a maximum of two e- with opposite spin

Question 32

Hund’s Rule

Answer 32

when there are orbitals of the same energy within a sublevel, electrons will half-full each orbital before pairing up –> this is cuz e- will repel each other and will want to be as far as possible

Question 33

Electron configuration for ions

Answer 33

cations (+) = e- are lost from the largest (outermost) shell –> highest n value, if n values are the same, then e- are removed from higher energy orbital first

anions (-)= e- are added to the nearest orbitals that are not yet full

Question 34

Paramagnetic VS diamagnetic

Answer 34

types of magnetism

PARAMAGNETIC
- unpaired electrons
- attracted to a magnetic field

DIAMAGNETIC
- paired electrons
- repelled by a magnetic field

Question 35

Identifying if an electron configuration is in its ground or excited state

Answer 35

• If valence e- are found in the order that we would predict, then the e- are in their ground state & as close to the nucleus as possible
• if the e- configuration is “out of order” (like if there is a gap), it means that e- are not in their ground state. –> they are in energy levels further from the nucleus than expected, they are in an excited state

Question 36

Exceptions to the electron configurations (Anomalous cases)

Answer 36

• some elements that do not behave according to their predicted e- configurations
• it turns out that when the “d” orbitals are half full or completely full, the e- can drop closer to the nucleus. –> This means that it is lower energy, or more stable, when the “d” orbitals are half or completely full
• these elements are: Nb, Mo, Cr, Ru, Rh, Pd, Ag, Pt, Au, Cu
• S^2 –> S^1
  –> only exception is Pd –> [Kr] 5s^2 4d^8 changes to [Kr]4d^10